Chemical formula for methane Methane Chemical formula for methane Methane is a colorless highly flammable gas that is less dense than air. The methane molecule consists of four hydrogen atoms covalently one carbon atom. The chemical formula for methane is CH4 which is the simplest hydrocarbon of the alkane series. The molar mass of methane is 16.043 g/mol Methane is a major component of natural gas and is also produced from by anaerobic bacteria in septic tanks and in the digestive systems of animals. Methane readily undergoes combustion producing carbon dioxide and water. Methane + oxygen ==> carbon dioxide + water + energy CH4 + 2O2 ==> CO2 + 2H2O + energy methane | Definition, meaning & more | Collins Dictionary methane for learners of EnglishPowered by COBUILD (miːθeɪn , US meθ- ) Definitions Methane is a colourless gas that has no smell . Natural gas consists mostly of methane. COBUILD Advanced English Dictionary. Copyright © HarperCollins Publishers methane (ˈmiːθeɪn ) Definitions noun a colourless odourless flammable gas , the simplest alkane and the main constituent of natural gas: used as a fuel .

Formula: CH4. → marsh gas , firedamp Collins English Dictionary. Copyright © HarperCollins Publishers Word origin of 'methane' methane in American English (ˈmɛθˌeɪn ; met hˈānˌ) Definitions noun a colorless, odorless, flammable, gaseous alkane, CH4, present in natural gas and formed by the decomposition of vegetable matter, as in marshes and mines , or produced artificially by heating carbon monoxide and hydrogen : it is the simplest alkane and is used as a fuel, a source of carbon black, etc. Webster’s New World College Dictionary, Fourth Edition, compiled by the editors of Webster’s New World Dictionaries. Copyright © 2010 by Houghton Mifflin Harcourt. Published by Houghton Mifflin Harcourt. All rights reserved. Word origin of 'methane' methane | chemical compound | Britannica.com chemical compound biogas Methane, colourless odourless gas that occurs abundantly in nature and as a product of certain human activities. Methane is the simplest member of the paraffin series of hydrocarbons and is among the most potent of the greenhouse gases .

Its chemical formula is CH4. Encyclopædia Britannica, Inc. Chemical properties of methane Methane is lighter than air , having a specific gravity of 0.554. It is only slightly soluble in water . It burns readily in air, forming carbon dioxide and water vapour; the flame is pale, slightly luminous, and very hot. The boiling point of methane is −162 °C (−259.6 °F) and the melting point is −182.5 °C (−296.5 °F). Methane in general is very stable, but mixtures of methane and air, with the methane content between 5 and 14 percent by volume, are explosive. Explosions of such mixtures have been frequent in coal mines and collieries and have been the cause of many mine disasters. The tetrahedral structure of methane (CH4) is explained in the VSEPR … Encyclopædia Britannica, Inc. Sources of methane In nature, methane is produced by the anaerobic bacterial decomposition of vegetable matter under water (where it is sometimes called marsh gas or swamp gas).

Wetlands are the major natural source of methane produced in this way. Other important natural sources of methane include termites (as a result of digestive processes), volcanoes , vents in the ocean floor, and methane hydrate deposits that occur along continental margins and beneath Antarctic ice and Arctic permafrost . Methane also is the chief constituent of natural gas , which contains from 50 to 90 percent methane (depending on the source), and occurs as a component of firedamp (flammable gas) along coal seams. Learn about the emission of methane, a greenhouse gas, by trees in wetland ecosystems. © Open University (A Britannica Publishing Partner) The tetrahedral geometry of methane: (A) stick-and-ball model and (B) showing bond angles and … Encyclopædia Britannica, Inc. wood The production and combustion of natural gas and coal are the major anthropogenic (human-associated) sources of methane. Activities such as the extraction and processing of natural gas and the destructive distillation of bituminous coal in the manufacture of coal gas and coke-oven gas result in the release of significant amounts of methane into the atmosphere .

Other human activities that are associated with methane production include biomass burning, livestock farming , and waste management (where bacteria produce methane as they decompose sludge in waste-treatment facilities and decaying matter in landfills). Uses of methane Methane is an important source of hydrogen and some organic chemicals. Methane reacts with steam at high temperatures to yield carbon monoxide and hydrogen; the latter is used in the manufacture of ammonia for fertilizers and explosives . Other valuable chemicals derived from methane include methanol , chloroform , carbon tetrachloride , and nitromethane. The incomplete combustion of methane yields carbon black , which is widely used as a reinforcing agent in rubber used for automobile tires . Role as a greenhouse gas Methane that is produced and released into the atmosphere is taken up by methane sinks, which include soil and the process of methane oxidation in the troposphere (the lowest atmospheric region). Most methane produced naturally is offset by its uptake into natural sinks. Anthropogenic methane production, however, can cause methane concentrations to increase more quickly than they are offset by sinks.

Increased concentrations of methane in the atmosphere contribute to the greenhouse effect , whereby greenhouse gases (particularly carbon dioxide , methane, and water vapour) absorb infrared radiation (net heat energy) and reradiate it back to Earth ’s surface, potentially trapping heat and producing substantial changes in climate . Increased atmospheric methane also adds to the greenhouse effect indirectly. For example, in methane oxidation hydroxyl radicals (OH−) remove methane by reacting with it to form carbon dioxide and water vapour, and as concentrations of atmospheric methane increase, concentrations of hydroxyl radicals decrease, effectively prolonging the atmospheric lifetime of methane. Learn More in these related articles: in climate (meteorology) Overview of Greenhouse Gases | Greenhouse Gas (GHG) Emissions | US EPA Overview of Greenhouse Gases Total Emissions in 2014 = 6,870 Million Metric Tons of CO2 equivalent . Larger image to save or print. Gases that trap heat in the atmosphere are called greenhouse gases. This section provides information on emissions and removals of the main greenhouse gases to and from the atmosphere. For more information on the science of climate change and other climate forcers, such as black carbon, please visit Climate Change Science .

6,870 million metric tons of CO2: What does that mean? An explanation of units: A million metric tons is equal to about 2.2 billion pounds, or 1 trillion grams. For comparison, a small car is likely to weigh a little more than 1 metric ton. Thus, a million metric tons is roughly the same mass as 1 million small cars! The U.S. Inventory uses metric units for consistency and comparability with other countries. For reference, a metric ton is a little bit larger (about 10%) than a U.S. "short" ton. GHG emissions are often measured in carbon dioxide (CO2) equivalent. To convert emissions of a gas into CO2 equivalent, its emissions are multiplied by the gas's Global Warming Potential (GWP) . The GWP takes into account the fact that many gases are more effective at warming Earth than CO2, per unit mass. The GWP values appearing in the Emissions webpages reflect the values used in the U.S. Inventory, which are drawn from the IPCC’s Fourth Assessment Report (AR4). For further discussion of GWPs and an estimate of GHG emissions using updated GWPs, see Annex 6 of the U.S.

Inventory and the IPCC's discussion on GWPs . Exit Carbon dioxide (CO2) : Carbon dioxide enters the atmosphere through burning fossil fuels (coal, natural gas, and oil), solid waste, trees and wood products, and also as a result of certain chemical reactions (e.g., manufacture of cement). Carbon dioxide is removed from the atmosphere (or "sequestered") when it is absorbed by plants as part of the biological carbon cycle. Methane (CH4) : Methane is emitted during the production and transport of coal, natural gas, and oil. Methane emissions also result from livestock and other agricultural practices and by the decay of organic waste in municipal solid waste landfills. Nitrous oxide (N2O) : Nitrous oxide is emitted during agricultural and industrial activities, as well as during combustion of fossil fuels and solid waste. Fluorinated gases : Hydrofluorocarbons, perfluorocarbons, sulfur hexafluoride, and nitrogen trifluoride are synthetic, powerful greenhouse gases that are emitted from a variety of industrial processes.

Fluorinated gases are sometimes used as substitutes for stratospheric ozone-depleting substances (e.g., chlorofluorocarbons, hydrochlorofluorocarbons, and halons). These gases are typically emitted in smaller quantities, but because they are potent greenhouse gases, they are sometimes referred to as High Global Warming Potential gases ("High GWP gases"). Each gas's effect on climate change depends on three main factors: How much of these gases are in the atmosphere? Concentration, or abundance, is the amount of a particular gas in the air. Larger emissions of greenhouse gases lead to higher concentrations in the atmosphere. Greenhouse gas concentrations are measured in parts per million, parts per billion, and even parts per trillion. One part per million is equivalent to one drop of water diluted into about 13 gallons of liquid (roughly the fuel tank of a compact car). To learn more about the increasing concentrations of greenhouse gases in the atmosphere, visit the Causes of Climate Change and the Climate Change Indicators: Atmospheric Concentrations of Greenhouse Gases pages. How long do they stay in the atmosphere?

Each of these gases can remain in the atmosphere for different amounts of time, ranging from a few years to thousands of years. All of these gases remain in the atmosphere long enough to become well mixed, meaning that the amount that is measured in the atmosphere is roughly the same all over the world, regardless of the source of the emissions. How strongly do they impact global temperatures? Some gases are more effective than others at making the planet warmer and "thickening the Earth's blanket." For each greenhouse gas, a Global Warming Potential (GWP) has been calculated to reflect how long it remains in the atmosphere, on average, and how strongly it absorbs energy. Gases with a higher GWP absorb more energy, per pound, than gases with a lower GWP, and thus contribute more to warming Earth. Global Warming Potential (100-year): 1 Carbon dioxide (CO2) is the primary greenhouse gas emitted through human activities. In 2014, CO2 accounted for about 80.9% of all U.S. greenhouse gas emissions from human activities.

Carbon dioxide is naturally present in the atmosphere as part of the Earth's carbon cycle (the natural circulation of carbon among the atmosphere, oceans, soil, plants, and animals). Human activities are altering the carbon cycle–both by adding more CO2 to the atmosphere and by influencing the ability of natural sinks, like forests, to remove CO2 from the atmosphere. While CO2 emissions come from a variety of natural sources, human-related emissions are responsible for the increase that has occurred in the atmosphere since the industrial revolution. 1 Note: All emission estimates from the Inventory of U.S. Greenhouse Gas Emissions and Sinks: 1990–2014 . Larger image to save or print. The main human activity that emits CO2 is the combustion of fossil fuels (coal, natural gas, and oil) for energy and transportation, although certain industrial processes and land-use changes also emit CO2. The main sources of CO2 emissions in the United States are described below. Electricity . Electricity is a significant source of energy in the United States and is used to power homes, business, and industry.

The combustion of fossil fuels to generate electricity is the largest single source of CO2 emissions in the nation, accounting for about 37 percent of total U.S. CO2 emissions and 30 percent of total U.S. greenhouse gas emissions in 2014. The type of fossil fuel used to generate electricity will emit different amounts of CO2. To produce a given amount of electricity, burning coal will produce more CO2 than oil or natural gas. Transportation . The combustion of fossil fuels such as gasoline and diesel to transport people and goods is the second largest source of CO2 emissions, accounting for about 31 percent of total U.S. CO2 emissions and 25 percent of total U.S. greenhouse gas emissions in 2014. This category includes transportation sources such as highway vehicles, air travel, marine transportation, and rail. Industry . Many industrial processes emit CO2 through fossil fuel combustion. Several processes also produce CO2 emissions through chemical reactions that do not involve combustion; for example, the production and consumption of mineral products such as cement, the production of metals such as iron and steel, and the production of chemicals.

Fossil fuel combustion from various industrial processes accounted for about 15 percent of total U.S. CO2 emissions and 12 percent of total U.S. greenhouse gas emissions in 2014. Note that many industrial processes also use electricity and therefore indirectly cause the emissions from the electricity production. Carbon dioxide is constantly being exchanged among the atmosphere, ocean, and land surface as it is both produced and absorbed by many microorganisms, plants, and animals. However, emissions and removal of CO2 by these natural processes tend to balance. Since the Industrial Revolution began around 1750, human activities have contributed substantially to climate change by adding CO2 and other heat-trapping gases to the atmosphere. In the United States, since 1990, the management of forests and non-agricultural land has acted as a net sink of CO2, which means that more CO2 is removed from the atmosphere, and stored in plants and trees, than is emitted. This sink offset about 11 percent of total emissions in 2014 and is discussed in more detail in the Land Use, Land-Use Change, and Forestry section.

To find out more about the role of CO2 warming the atmosphere and its sources, visit the Causes of Climate Change page and the Climate Change Indicators page in the Science section. Emissions and Trends Carbon dioxide (CO2) emissions in the United States increased by about 9 percent between 1990 and 2014. Since the combustion of fossil fuel is the largest source of greenhouse gas emissions in the United States, changes in emissions from fossil fuel combustion have historically been the dominant factor affecting total U.S. emission trends. Changes in CO2 emissions from fossil fuel combustion are influenced by many long-term and short-term factors, including population growth, economic growth, changing energy prices, new technologies, changing behavior, and seasonal temperatures. Between 1990 and 2014, the increase in CO2 emissions corresponded with increased energy use by an expanding economy and population, and an overall growth in emissions from electricity generation. Transportation emissions also contributed to the 9 percent increase, largely due to an increase in miles traveled by motor vehicles. Larger image to save or print. Reducing Carbon Dioxide Emissions The most effective way to reduce carbon dioxide (CO2) emissions is to reduce fossil fuel consumption.

Many strategies for reducing CO2 emissions from energy are cross-cutting and apply to homes, businesses, industry, and transportation. EPA is taking common sense regulatory actions to reduce greenhouse gas emissions from our nation's largest sources, including power plants and motor vehicles. Examples of How Emissions Can be Reduced Energy Efficiency Improving the insulation of buildings, traveling in more fuel-efficient vehicles, and using more efficient electrical appliances are all ways to reduce energy consumption, and thus CO2 emissions. See EPA's ENERGY STAR® program for more information on energy-efficient appliances. See EPA's and DOE's fueleconomy.gov site for more information on fuel-efficient vehicles. Learn about EPA's motor vehicle standards that improve vehicle efficiency and save drivers money. Energy Conservation Reducing personal energy use by turning off lights and electronics when not in use reduces electricity demand. Reducing distance traveled in vehicles reduces petroleum consumption. Both are ways to reduce energy CO2 emissions through conservation. Learn more about What You Can Do at Home , at School , in the Office , and on the Road to save energy and reduce your carbon footprint. Fuel Switching Producing more energy from renewable sources and using fuels with lower carbon contents are ways to reduce carbon emissions.

Carbon Capture and Sequestration Carbon dioxide capture and sequestration is a set of technologies that can potentially greatly reduce CO2 emissions from new and existing coal- and gas-fired power plants, industrial processes, and other stationary sources of CO2. Learn more . * Carbon dioxide's lifetime is poorly defined because the gas is not destroyed over time, but instead moves among different parts of the ocean-atmosphere-land system. Some of the excess carbon dioxide will be absorbed quickly (for example, by the ocean surface), but some will remain in the atmosphere for thousands of years, due in part to the very slow process by which carbon is transferred to ocean sediments. References Global Warming Potential (100-year): 25 Methane (CH4) is the second most prevalent greenhouse gas emitted in the United States from human activities. In 2014, CH4 accounted for about 11 percent of all U.S. greenhouse gas emissions from human activities. Methane is emitted by natural sources such as wetlands, as well as human activities such as leakage from natural gas systems and the raising of livestock. Natural processes in soil and chemical reactions in the atmosphere help remove CH4 from the atmosphere.

Methane's lifetime in the atmosphere is much shorter than carbon dioxide (CO2), but CH4 is more efficient at trapping radiation than CO2. Pound for pound, the comparative impact of CH4 on climate change is more than 25 times greater than CO2 over a 100-year period. Globally, over 60 percent of total CH4 emissions come from human activities. 1 Methane is emitted from industry, agriculture, and waste management activities, described below. Note: All emission estimates from the Inventory of U.S. Greenhouse Gas Emissions and Sinks: 1990–2014 . Larger image to save or print. Industry . Natural gas and petroleum systems are the largest source of CH4 emissions from industry in the United States. Methane is the primary component of natural gas. Some CH4 is emitted to the atmosphere during the production, processing, storage, transmission, and distribution of natural gas. Because gas is often found alongside petroleum, the production, refinement, transportation, and storage of crude oil is also a source of CH4 emissions. For more information, see the Inventory of U.S. Greenhouse Gas Emissions and Sinks sections on Natural Gas Systems and Petroleum Systems.

Agriculture . Domestic livestock such as cattle, buffalo, sheep, goats, and camels produce large amounts of CH4 as part of their normal digestive process. Also, when animals' manure is stored or managed in lagoons or holding tanks, CH4 is produced. Because humans raise these animals for food, the emissions are considered human-related. Globally, the Agriculture sector is the primary source of CH4 emissions. For more information, see the Inventory of U.S. Greenhouse Gas Emissions and Sinks Agriculture chapter. Waste from Homes and Businesses . Methane is generated in landfills as waste decomposes and in the treatment of wastewater. Landfills are the third largest source of CH4 emissions in the United States. For more information see the U.S. Inventory's Waste chapter. Methane is also emitted from a number of natural sources. Wetlands are the largest source, emitting CH4 from bacteria that decompose organic materials in the absence of oxygen. Smaller sources include termites, oceans, sediments, volcanoes, and wildfires.

To find out more about the role of CH4 in warming the atmosphere, and its sources, visit the Causes of Climate Change page and the Climate Change Indicators page in the Science section. Emissions and Trends Methane (CH4) emissions in the United States decreased by 6 percent between 1990 and 2014. During this time period, emissions increased from sources associated with agricultural activities, while emissions decreased from sources associated with the exploration and production of natural gas and petroleum products. Note: All emission estimates from the Inventory of U.S. Greenhouse Gas Emissions and Sinks: 1990-2014 . These estimates use a global warming potential for methane of 25, based on reporting requirements under the United Nations Framework Convention on Climate Change. Larger image to save or print. Reducing Methane Emissions There are a number of ways to reduce methane (CH4) emissions. Some examples are discussed below. EPA has a series of voluntary programs for reducing CH4 emissions, and is supporting the president's Strategy to Reduce Methane Emissions (15 pp, 1.9 M, About PDF ).

EPA also supports the Global Methane Initiative Exit , an international partnership encouraging global methane reduction strategies. Examples of Reduction Opportunities for Methane Emissions Source How Emissions Can be Reduced Industry Upgrading the equipment used to produce, store, and transport oil and gas can reduce many of the leaks that contribute to CH4 emissions. Methane from coal mines can also be captured and used for energy. Learn more about the EPA's Natural Gas STAR Program and Coalbed Methane Outreach Program . Agriculture Methane can be reduced and captured by altering manure management strategies at livestock operations or animal feeding practices. Learn more about these strategies and EPA's AgSTAR Program . Waste from Homes and Businesses Because CH4 emissions from landfill gas are a major source of CH4 emissions in the United States, emission controls that capture landfill CH4 are an effective reduction strategy. Learn more about these opportunities and the EPA's Landfill Methane Outreach Program . References Global Warming Potential (100-year): 298 In 2014, nitrous oxide (N2O) accounted for about 6 percent of all U.S. greenhouse gas emissions from human activities.

Nitrous oxide is naturally present in the atmosphere as part of the Earth's nitrogen cycle, and has a variety of natural sources. However, human activities such as agriculture, fossil fuel combustion, wastewater management, and industrial processes are increasing the amount of N2O in the atmosphere. Nitrous oxide molecules stay in the atmosphere for an average of 114 years before being removed by a sink or destroyed through chemical reactions. The impact of 1 pound of N2O on warming the atmosphere is almost 300 times that of 1 pound of carbon dioxide. Globally, about 40 percent of total N2O emissions come from human activities. 1 Nitrous oxide is emitted from agriculture, transportation, and industry activities, described below. Note: All emission estimates from the Inventory of U.S. Greenhouse Gas Emissions and Sinks: 1990–2014 . Larger image to save or print. Agriculture . Nitrous oxide is emitted when people add nitrogen to the soil through the use of synthetic fertilizers. Agricultural soil management is the largest source of N2O emissions in the United States, accounting for about 79 percent of total U.S.

N2O emissions in 2014. Nitrous oxide is also emitted during the breakdown of nitrogen in livestock manure and urine, which contributed to 4 percent of N2O emissions in 2014. Transportation . Nitrous oxide is emitted when transportation fuels are burned. Motor vehicles, including passenger cars and trucks, are the primary source of N2O emissions from transportation. The amount of N2O emitted from transportation depends on the type of fuel and vehicle technology, maintenance, and operating practices. Industry . Nitrous oxide is generated as a byproduct during the production of nitric acid, which is used to make synthetic commercial fertilizer, and in the production of adipic acid, which is used to make fibers, like nylon, and other synthetic products. Nitrous oxide emissions occur naturally through many sources associated with the nitrogen cycle, which is the natural circulation of nitrogen among the atmosphere, plants, animals, and microorganisms that live in soil and water. Nitrogen takes on a variety of chemical forms throughout the nitrogen cycle, including N2O. Natural emissions of N2O are mainly from bacteria breaking down nitrogen in soils and the oceans.

Nitrous oxide is removed from the atmosphere when it is absorbed by certain types of bacteria or destroyed by ultraviolet radiation or chemical reactions. To find out more about the role of N2O in warming the atmosphere and its sources, visit the Causes of Climate Change page and the Climate Change Indicators page in the Science section. Emissions and Trends Nitrous oxide (N2O) emissions in the United States have decreased by 1 percent between 1990 and 2014. This decrease in emissions is due in part to a decrease in emissions from mobile combustion resulting from emission control standards for on-road vehicles. Nitrous oxide emissions from agricultural soils have varied during this period and were about 5 percent higher in 2014 than in 1990. Going forward, N2O emissions are projected to increase by 5 percent between 2005 and 2020, driven largely by increases in emissions from agricultural activities. 2 There are a number of ways to reduce emissions of nitrous oxide (N2O), discussed below.

Examples of Reduction Opportunities for Nitrous Oxide Emissions Emissions Source Examples of How Emissions Can be Reduced Agriculture The application of fertilizers accounts for the majority of N2O emissions. Emissions can be reduced by reducing nitrogen-based fertilizer applications and applying fertilizers more efficiently, 3 as well as following better manure management practices. Transportation Nitrous oxide is a byproduct of fuel combustion, so reducing mobile fuel consumption in motor vehicles can reduce transportation emissions. Additionally, the introduction of pollution control technologies, such as catalytic converters to reduce exhaust pollutants from passenger cars, can also reduce emissions of N2O. Industry Nitrous oxide is generally emitted from industry through fossil fuel combustion, so technological upgrades and fuel switching are effective ways to reduce industry emissions of N2O. Production of adipic acid results in N2O emissions that can be reduced through technological upgrades. References NF3: 17,200 SF6: 22,800 Unlike many other greenhouse gases, fluorinated gases have no natural sources and only come from human-related activities. They are emitted through a variety of industrial processes such as aluminum and semiconductor manufacturing.

Many fluorinated gases have very high global warming potentials (GWPs) relative to other greenhouse gases, so small atmospheric concentrations can have large effects on global temperatures. They can also have long atmospheric lifetimes—in some cases, lasting thousands of years. Like other long-lived greenhouse gases, fluorinated gases are well-mixed in the atmosphere, spreading around the world after they are emitted. Fluorinated gases are removed from the atmosphere only when they are destroyed by sunlight in the far upper atmosphere. In general, fluorinated gases are the most potent and longest lasting type of greenhouse gases emitted by human activities. There are four main categories of fluorinated gases—hydrofluorocarbons (HFCs), perfluorocarbons (PFCs), sulfur hexafluoride (SF6), and nitrogen trifluoride (NF3). The largest sources of fluorinated gas emissions are described below. Larger image to save or print. Substitution for Ozone-Depleting Substances. Hydrofluorocarbons are used as refrigerants, aerosol propellants, solvents, and fire retardants.

The major emissions source of these compounds is their use as refrigerants—for example, in air conditioning systems in both vehicles and buildings. These chemicals were developed as a replacement for chlorofluorocarbons (CFCs) and hydrochlorofluorocarbons (HCFCs) because they do not deplete the stratospheric ozone layer. Chlorofluorocarbons and HCFCs are being phased out under an international agreement, called the Montreal Protocol. Unfortunately, HFCs are potent greenhouse gases with long atmospheric lifetimes and high GWPs, and they are released into the atmosphere through leaks, servicing, and disposal of equipment in which they are used. Industry. Perfluorocarbons are compounds produced as a byproduct of various industrial processes associated with aluminum production and the manufacturing of semiconductors. Like HFCs, PFCs generally have long atmospheric lifetimes and high GWPs. Sulfur hexafluoride is used in magnesium processing and semiconductor manufacturing, as well as a tracer gas for leak detection. HFC-23 is produced as a byproduct of HCFC-22 production. Transmission and Distribution of Electricity.

Sulfur hexafluoride is used in electrical transmission equipment, including circuit breakers. The GWP of SF6 is 22,800, making it the most potent greenhouse gas that the Intergovernmental Panel on Climate Change has evaluated. To find out more about the role of fluorinated gases in warming the atmosphere, and their sources, visit the Causes of Climate Change page and the Climate Change Indicators page in the Science section. Emissions and Trends Overall, fluorinated gas emissions in the United States have increased by about 77 percent between 1990 and 2014. This increase has been driven by a 258 percent increase in emissions of hydrofluorocarbons (HFCs) since 1990, as they have been widely used as a substitute for ozone-depleting substances. Emissions of perfluorocarbons (PFCs) and sulfur hexafluoride (SF6) have actually declined during this time due to emission reduction efforts in the aluminum production industry (PFCs) and the electricity transmission and distribution industry (SF6).

Going forward, HFC emissions are projected to grow by nearly 141 percent between 2005 and 2020 as demands for refrigeration continue to grow and as more ozone-depleting substances are replaced. During this same period, emissions of SF6 are expected to decline by over 25 percent, while emissions of PFCs are projected to remain flat. Since the emissions from HFCs are far greater than the emissions of the other gases on a carbon dioxide-equivalence basis, substantial growth in emissions of fluorinated gases is expected to continue. 1 Larger image to save or print. Reducing Fluorinated Gas Emissions Because most fluorinated gases have a very long atmospheric lifetime, it will take many years to see a noticeable decline in current concentrations. However, there are a number of ways to reduce emissions of fluorinated gases, described below. Examples of Reduction Opportunities for Fluorinated Gases Emissions Source The PFC Reduction/Climate Partnership for the Semiconductor Industry Electricity Transmission and Distribution Sulfur hexafluoride is an extremely potent greenhouse gas that is used for several purposes when transmitting electricity through the power grid.

EPA is working with industry to reduce emissions through the SF6 Emission Reduction Partnership for Electric Power Systems , which promotes leak detection and repair, use of recycling equipment, and employee training. Transportation Hydrofluorocarbons (HFCs) are released through the leakage of refrigerants used in vehicle air-conditioning systems. Leakage can be reduced through better system components, and through the use of alternative refrigerants with lower global warming potentials than those presently used. One important way the EPA is working to reduce HFC emissions is through its light-duty and heavy-duty vehicle standards . References (10 pp, 577 K, About PDF ). U.S. Department of State, Washington, DC, USA. Multimedia: Balanced Equation | Chapter 6, Lesson 1 | Middle School Chemistry Middle School Chemistry Chapter 6, Lesson 1 Multimedia Methane and Oxygen React In the reaction, the the bonds in the methane and oxygen come apart, the atoms rearrange and then re-bond to form water and carbon dioxide. The little number written at the lower right after an atom (subscript) tells how many of that atom are in the molecule.

The big number written in front of a molecule (coefficient) shows how many of that molecule there are. All the atoms in the products come from the atoms in the reactants. Combustion of Methane What Is the Chemical Composition of Farts? What Is the Chemical Composition of Farts? What Is the Chemical Composition of Farts? What Flatus or Farts Are Made Of The chemical composition of a fart depends on a person's body chemistry, but typically includes stinky sulfur compounds. Digital Vision, Getty Images Updated August 07, 2015. Question: What Is the Chemical Composition of Farts? Farts are the common name for flatus or flatulence. Have you ever wondered what farts are made of and whether they are the same for everyone? Here's a look at the chemical composition of farts. Answer: Have you ever wondered about the chemical composition of farts, or to be more technical, flatus? The exact chemical composition of human flatulence varies from one person to another, based on his or her biochemistry, the bacteria inhabiting the colon and the foods that were eaten. If the gas results from ingesting air, the chemical composition will approximate that of air .

If the fart arises from digestion or bacterial production, the chemistry may be more exotic. Farts consist primarily of nitrogen, the principal gas in air, along with a significant amount of carbon dioxide . A typical breakdown of the chemical composition of farts is: Nitrogen: 20-90% continue reading below our video 10 Facts About the Titanic That You Don't Know Lighting Farts on Fire - The Blue Flame Human flatus may contain hydrogen gas and/or methane, which are flammable. If sufficient amounts of these gases are present, it's possible to light the fart on fire . Keep in mind, not all farts are flammable. Although flatus has great YouTube fame for producing a blue flame, it turns out only about half of people have the archaea (bacteria) in their bodies that are necessary to produce methane. If you don't make methane, you may still be able to ignite your farts (a dangerous practice!), but the flame will be yellow or possibly orange rather than blue. The Smell of Farts Flatus often stinks!

There are several chemicals that contribute to the smell of farts: skatole (by-product of meat digestion) indole (by-product of meat digestion) methanethiol (a sulfur compound) Natural Gas Basics » Gulf Coast Midwest Energy Natural Gas Basics Natural Gas Basics Natural gas and crude oil are called hydrocarbons because both are composed of carbon and hydrogen atoms. Natural gas molecules are generally shorter; four carbon atoms or less. Crude oil molecules contain five or more carbon atoms per molecule. Natural gas is both a clean burning fuel source and the primary feedstock into the U.S. petrochemical industry. Raw natural gas in fields is processed to separate the methane out from butane, propane, and large amounts of ethane. The three heavier liquids are prone to condensation in natural gas pipelines. Natural Gas Liquids and Petrochemicals Methane - One carbon atom; chemical formula CH4. The principal use of methane is as a fuel. The natural gas that is delivered to your home is almost pure methane. Methane is also an upstream component of many important industrial chemicals: Methanol - A primary raw material in automotive windshield wash and also used as a racing fuel.

Chemical derivatives of methanol also play an important role in our everyday lives: Ethane - Two carbon atoms; chemical formula C2H6. Ethane is the second-largest component of natural gas, typically comprising up to 6% of the volume of gas produced from a natural gas field. Ethylene - The vast majority of ethane is consumed in the production of ethylene by steam cracking. Consumer products made from ethylene derivatives represent a large part of the petrochemical industry here in the U.S.: Propane - Three carbon atoms; chemical formula C3H8. Truck fleets, forklifts, barbecue grills, portable stoves and even the new Roush Ford F-150 pick-up use propane fuel. Propane's octane rating is noticeably higher than gasoline at 110. Many homes not connected to municipal (methane) gas pipelines use propane for their appliances and furnaces. Other fuel uses for propane: Refrigerators. Molecular Structure & Bonding 2 180º In the three examples shown above, the central atom (carbon) does not have any non-bonding valence electrons; consequently the configuration may be estimated from the number of bonding partners alone.

For molecules of water and ammonia, however, the non-bonding electrons must be included in the calculation. In each case there are four regions of electron density associated with the valence shell so that a tetrahedral bond angle is expected. The measured bond angles of these compounds (H2O 104.5º & NH3 107.3º) show that they are closer to being tetrahedral than trigonal or linear. Of course, it is the configuration of atoms (not electrons) that defines the the shape of a molecule, and in this sense ammonia is said to be pyramidal (not tetrahedral). The compound boron trifluoride, BF3, does not have non-bonding valence electrons and the configuration of its atoms is trigonal. Nice treatments of VSEPR theory have been provided by Oxford and Purdue . Click on the university name to visit their site. The best way to study the three-dimensional shapes of molecules is by using molecular models. Many kinds of model kits are available to students and professional chemists. Some of the useful features of physical models can be approximated by the model viewing applet Jmol.

This powerful visualization tool allows the user to move a molecular stucture in any way desired. Atom distances and angles are easily determined. To measure a distance, double-click on two atoms. To measure a bond angle, do a double-click, single-click, double-click on three atoms. To measure a torsion angle, do a double-click, single-click, single-click, double-click on four atoms. A pop-up menu of commands may be accessed by the right button on a PC or a control-click on a Mac while the cursor is inside the display frame. You may examine several Jmol models of compounds discussed above by . One way in which the shapes of molecules manifest themselves experimentally is through molecular dipole moments. A molecule which has one or more polar covalent bonds may have a dipole moment as a result of the accumulated bond dipoles. In the case of water, we know that the O-H covalent bond is polar, due to the different electronegativities of hydrogen and oxygen. Since there are two O-H bonds in water, their bond dipoles will interact and may result in a molecular dipole which can be measured. The following diagram shows four possible orientations of the O-H bonds.

The bond dipoles are colored magenta and the resulting molecular dipole is colored blue. In the linear configuration (bond angle 180º) the bond dipoles cancel, and the molecular dipole is zero. For other bond angles (120 to 90º) the molecular dipole would vary in size, being largest for the 90º configuration. In a similar manner the configurations of methane (CH4) and carbon dioxide (CO2) may be deduced from their zero molecular dipole moments. Since the bond dipoles have canceled, the configurations of these molecules must be tetrahedral (or square-planar) and linear respectively. The case of methane provides insight to other arguments that have been used to confirm its tetrahedral configuration. For purposes of discussion we shall consider three other configurations for CH4, square-planar, square-pyramidal and triangular-pyramidal. Models of these possibilities may be examined by . Substitution of one hydrogen by a chlorine atom gives a CH3Cl compound. Since the tetrahedral, square-planar and square-pyramidal configurations have structurally equivalent hydrogen atoms, they would each give a single substitution product.

However, in the trigonal-pyramidal configuration one hydrogen (the apex) is structurally different from the other three (the pyramid base). Substitution in this case should give two different CH3Cl compounds if all the hydrogens react. In the case of disubstitution, the tetrahedral configuration of methane would lead to a single CH2Cl2 product, but the other configurations would give two different CH2Cl2 compounds. These substitution possibilities are shown in the above insert. Isomers For an example Click Here . Distinguishing Carbon Atoms When discussing structural formulas, it is often useful to distinguish different groups of carbon atoms by their structural characteristics. A primary carbon (1º) is one that is bonded to no more than one other carbon atom. A secondary carbon (2º) is bonded to two other carbon atoms, and tertiary (3º) and quaternary (4º) carbon atoms are bonded respectively to three and four other carbons. The three C5H12 isomers shown below illustrate these terms. Structural differences may occur within these four groups, depending on the molecular constitution.

In the formula on the right all four 1º-carbons are structurally equivalent (remember the tetrahedral configuration of tetravalent carbon); however the central formula has two equivalent 1º-carbons (bonded to the 3º carbon on the left end) and a single, structurally different 1º-carbon (bonded to the 2º-carbon) at the right end. Similarly, the left-most formula has two structurally equivalent 2º-carbons (next to the ends of the chain), and a structurally different 2º-carbon in the middle of the chain. A consideration of molecular symmetry helps to distinguish structurally equivalent from nonequivalent atoms and groups. The ability to distinguish structural differences of this kind is an essential part of mastering organic chemistry. It will come with practice and experience. Our ability to draw structural formulas for molecules is remarkable. To see how this is done Click Here . Formula Analysis Analysis of Molecular Formulas Although structural formulas are essential to the unique description of organic compounds, it is interesting and instructive to evaluate the information that may be obtained from a molecular formula alone.

Three useful rules may be listed: The number of hydrogen atoms that can be bonded to a given number of carbon atoms is limited by the valence of carbon. For compounds of carbon and hydrogen (hydrocarbons) the maximum number of hydrogen atoms that can be bonded to n carbons is 2n + 2 (n is an integer). In the case of methane, CH4, n=1 & 2n + 2 = 4. The origin of this formula is evident by considering a hydrocarbon made up of a chain of carbon atoms. Here the middle carbons will each have two hydrogens and the two end carbons have three hydrogens each. Thus, a six-carbon chain (n = 6) may be written H-(CH2)6-H, and the total hydrogen count is (2 x 6) + 2 = 14. The presence of oxygen (valence = 2) does not change this relationship, so the previously described C4H10O isomers follow the rule, n=4 & 2n + 2 = 10.

Halogen atoms (valence = 1) should be counted equivalent to hydrogen, as illustrated by C3H5Cl3, n = 3 & 2n + 2 = 8 = (5 + 3). If nitrogen is present, each nitrogen atom (valence = 3) increases the maximum number of hydrogens by one. Some Plausible Molecular Formulas C5H9O2, C4H5ClBr, C6H11N2O, C10H18NCl2 The number of hydrogen atoms in stable compounds of carbon, hydrogen & oxygen reflects the number of double bonds and rings in their structural formulas. Consider a hydrocarbon with a molecular structure consisting of a simple chain of four carbon atoms, CH3CH2CH2CH3. The molecular formula is C4H10 (the maximum number of bonded hydrogens by the 2n + 2 rule). If the four carbon atoms form a ring, two hydrogens must be lost. Similarly, the introduction of a double bond entails the loss of two hydrogens, and a triple bond the loss of four hydrogens.

From the above discussion and examples it should be clear that the molecular formula of a hydrocarbon (CnHm) provides information about the number of rings and/or double bonds that must be present in its structural formula. By rule #2 m must be an even number, so if m < (2n + 2) the difference is also an even number that reflects any rings and double bonds. A triple bond is counted as two double bonds. The presence of one or more nitrogen atoms or halogen substituents requires a modified analysis. The above formula may be extended to such compounds by a few simple principles: The presence of oxygen does not alter the relationship. All halogens present in the molecular formula must be replaced by hydrogen. Each nitrogen in the formula must be replaced by a CH moiety. Resonance Resonance Kekulé structural formulas are essential tools for understanding organic chemistry. However, the structures of some compounds and ions cannot be represented by a single formula. For example, sulfur dioxide (SO2) and nitric acid (HNO3) may each be described by two equivalent formulas (equations 1 & 2). For clarity the two ambiguous bonds to oxygen are given different colors in these formulas.

1) sulfur dioxide 2) nitric acid If only one formula for sulfur dioxide was correct and accurate, then the double bond to oxygen would be shorter and stronger than the single bond. Since experimental evidence indicates that this molecule is bent (bond angle 120º) and has equal length sulfur : oxygen bonds (1.432 Å), a single formula is inadequate, and the actual structure resembles an average of the two formulas. This averaging of electron distribution over two or more hypothetical contributing structures (canonical forms) to produce a hybrid electronic structure is called resonance. Likewise, the structure of nitric acid is best described as a resonance hybrid of two structures, the double headed arrow being the unique symbol for resonance. The above examples represent one extreme in the application of resonance. Here, two structurally and energetically equivalent electronic structures for a stable compound can be written, but no single structure provides an accurate or even an adequate representation of the true molecule. In cases such as these, the electron delocalization described by resonance enhances the stability of the molecules, and compounds or ions composed of such molecules often show exceptional stability. 3) formaldehyde The electronic structures of most covalent compounds do not suffer the inadequacy noted above.

Thus, completely satisfactory Kekulé formulas may be drawn for water (H2O), methane (CH4) and acetylene C2H2). Nevertheless, the principles of resonance are very useful in rationalizing the chemical behavior of many such compounds. For example, the carbonyl group of formaldehyde (the carbon-oxygen double bond) reacts readily to give addition products. The course of these reactions can be explained by a small contribution of a dipolar resonance contributor, as shown in equation 3. Here, the first contributor (on the left) is clearly the best representation of this molecular unit, since there is no charge separation and both the carbon and oxygen atoms have achieved valence shell neon-like configurations by covalent electron sharing. If the double bond is broken heterolytically, formal charge pairs result, as shown in the other two structures. The preferred charge distribution will have the positive charge on the less electronegative atom (carbon) and the negative charge on the more electronegative atom (oxygen). Therefore the middle formula represents a more reasonable and stable structure than the one on the right. The application of resonance to this case requires a weighted averaging of these canonical structures.

The double bonded structure is regarded as the major contributor, the middle structure a minor contributor and the right hand structure a non-contributor. Since the middle, charge-separated contributor has an electron deficient carbon atom, this explains the tendency of electron donors (nucleophiles) to bond at this site. The basic principles of the resonance method may now be summarized. For a given compound, a set of Lewis / Kekulé structures are written, keeping the relative positions of all the component atoms the same. These are the canonical forms to be considered, and all must have the same number of paired and unpaired electrons. The following factors are important in evaluating the contribution each of these canonical structures makes to the actual molecule. The number of covalent bonds in a structure. (The greater the bonding, the more important and stable the contributing structure.) Formal charge separation. (Other factors aside, charge separation decreases the stability and importance of the contributing structure.) Electronegativity of charge bearing atoms and charge density. (High charge density is destabilizing. Positive charge is best accommodated on atoms of low electronegativity, and negative charge on high electronegative atoms.) The stability of a resonance hybrid is always greater than the stability of any canonical contributor.

Consequently, if one canonical form has a much greater stability than all others, the hybrid will closely resemble it electronically and energetically. This is the case for the carbonyl group (eq.3). The left hand C=O structure has much greater total bonding than either charge-separated structure, so it describes this functional group rather well. On the other hand, if two or more canonical forms have identical low energy structures, the resonance hybrid will have exceptional stabilization and unique properties. This is the case for sulfur dioxide (eq.1) and nitric acid (eq.2). 4) carbon monoxide 5) azide anion To illustrate these principles we shall consider carbon monoxide (eq.4) and azide anion (eq.5). In each case the most stable canonical form is on the left. For carbon monoxide, the additional bonding is more important than charge separation. Furthermore, the double bonded structure has an electron deficient carbon atom (valence shell sextet). A similar destabilizing factor is present in the two azide canonical forms on the top row of the bracket (three bonds vs. four bonds in the left most structure).

The bottom row pair of structures have four bonds, but are destabilized by the high charge density on a single nitrogen atom. All the examples on this page demonstrate an important restriction that must be remembered when using resonance. No atoms change their positions within the common structural framework. Only electrons are moved. Orbitals Atomic and Molecular Orbitals A more detailed model of covalent bonding requires a consideration of valence shell atomic orbitals. For second period elements such as carbon, nitrogen and oxygen, these orbitals have been designated 2s, 2px, 2py & 2pz. The spatial distribution of electrons occupying each of these orbitals is shown in the diagram below. Very nice displays of orbitals may be found at the following sites: M. Winter, Sheffield University The valence shell electron configuration of carbon is 2s2, 2px1, 2py1 & 2pz0. If this were the configuration used in covalent bonding, carbon would only be able to form two bonds. In this case, the valence shell would have six electrons- two shy of an octet.

However, the tetrahedral structures of methane and carbon tetrachloride demonstrate that carbon can form four equivalent bonds, leading to the desired octet. In order to explain this covalent bonding, Linus Pauling proposed an orbital hybridization model in which all the valence shell electrons of carbon are reorganized. Hybrid Orbitals In order to explain the structure of methane (CH4), the 2s and three 2p orbitals are converted to four equivalent hybrid atomic orbitals, each having 25% s and 75% p character, and designated sp3. These hybrid orbitals have a specific orientation, and the four are naturally oriented in a tetrahedral fashion. Thus, the four covalent bonds of methane consist of shared electron pairs with four hydrogen atoms in a tetrahedral configuration, as predicted by VSEPR theory. Molecular Orbitals Just as the valence electrons of atoms occupy atomic orbitals (AO), the shared electron pairs of covalently bonded atoms may be thought of as occupying molecular orbitals (MO). It is convenient to approximate molecular orbitals by combining or mixing two or more atomic orbitals. In general, this mixing of n atomic orbitals always generates n molecular orbitals.

The hydrogen molecule provides a simple example of MO formation. In the following diagram, two 1s atomic orbitals combine to give a sigma (σ) bonding (low energy) molecular orbital and a second higher energy MO referred to as an antibonding orbital. The bonding MO is occupied by two electrons of opposite spin, the result being a covalent bond. The notation used for molecular orbitals parallels that used for atomic orbitals. Thus, s-orbitals have a spherical symmetry surrounding a single nucleus, whereas σ-orbitals have a cylindrical symmetry and encompass two (or more) nuclei. In the case of bonds between second period elements, p-orbitals or hybrid atomic orbitals having p-orbital character are used to form molecular orbitals. For example, the sigma molecular orbital that serves to bond two fluorine atoms together is generated by the overlap of p-orbitals (part A below), and two sp3 hybrid orbitals of carbon may combine to give a similar sigma orbital. When these bonding orbitals are occupied by a pair of electrons, a covalent bond, the sigma bond results.

Although we have ignored the remaining p-orbitals, their inclusion in a molecular orbital treatment does not lead to any additional bonding, as may be shown by activating the fluorine correlation diagram below. Another type of MO (the π orbital) may be formed from two p-orbitals by a lateral overlap, as shown in part A of the following diagram. Since bonds consisting of occupied π-orbitals (pi-bonds) are weaker than sigma bonds, pi-bonding between two atoms occurs only when a sigma bond has already been established. Thus, pi-bonding is generally found only as a component of double and triple covalent bonds. Since carbon atoms involved in double bonds have only three bonding partners, they require only three hybrid orbitals to contribute to three sigma bonds. A mixing of the 2s-orbital with two of the 2p orbitals gives three sp2 hybrid orbitals, leaving one of the p-orbitals unused. Two sp2 hybridized carbon atoms are then joined together by sigma and pi-bonds (a double bond), as shown in part B. The manner in which atomic orbitals overlap to form molecular orbitals is actually more complex than the localized examples given above.

These are useful models for explaining the structure and reactivity of many organic compounds, but modern molecular orbital theory involves the creation of an orbital correlation diagram. Two examples of such diagrams for the simple diatomic elements F2 and N2 will be drawn above when the appropriate button is clicked. The 1s and 2s atomic orbitals do not provide any overall bonding, since orbital overlap is minimal, and the resulting sigma bonding and antibonding components would cancel. In both these cases three 2p atomic orbitals combine to form a sigma and two pi-molecular orbitals, each as a bonding and antibonding pair. The overall bonding order depends on the number of antibonding orbitals that are occupied. The subtle change in the energy of the σ2p bonding orbital, relative to the two degenerate π-bonding orbitals, is due to s-p hybridization that is unimportant to the present discussion. One example of the advantage offered by the molecular orbital approach to bonding is the oxygen molecule . Here, the correlation diagram correctly accounts for the paramagnetic character of this simple diatomic compound.

Likewise, the orbital correlation diagram for methane provides another example of the difference in electron density predicted by molecular orbital calculations from that of the localized bond model. Click on the compound names for these displays. A cartoon of the p and π orbitals of a double bond may be examined by . A model of the π orbitals of ethene may be examined by . The p-orbitals in these model are represented by red and blue colored spheres or ellipses, which represent different phases, defined by the mathematical wave equations for such orbitals. Finally, in the case of carbon atoms with only two bonding partners only two hybrid orbitals are needed for the sigma bonds, and these sp hybrid orbitals are directed 180º from each other. Two p-orbitals remain unused on each sp hybridized atom, and these overlap to give two pi-bonds following the formation of a sigma bond (a triple bond), as shown below. The various hybridization states of carbon may be examined by . Practice Problems The following problems explore many of the concepts discussed above. They include recognition of different molecular shapes, obtaining structural information from molecular formulas, evaluation of line and condensed structural formulas, formal charges, hybridization and resonance.

This page is the property of William Reusch. Comments, questions and errors should be sent to whreusch@msu.edu . These pages are provided to the IOCD to assist in capacity building in chemical education. 05/05/2013

Question: Which gas has the chemical formula CH4?
Answer: METHANE